1 / 28

REDOX

REDOX. Oxidation-Reduction Reactions. Oxidation-Reduction Reactions. Electrons are transferred from 1 atom to another. All single-replacement & combustion rxns are redox rxns. Oxidation. = loss of electrons. L OSS of E LECTRONS = O XIDATION LEO. Reduction. = gain of electrons.

aaron-velazquez
Download Presentation

REDOX

An Image/Link below is provided (as is) to download presentation Download Policy: Content on the Website is provided to you AS IS for your information and personal use and may not be sold / licensed / shared on other websites without getting consent from its author. Content is provided to you AS IS for your information and personal use only. Download presentation by click this link. While downloading, if for some reason you are not able to download a presentation, the publisher may have deleted the file from their server. During download, if you can't get a presentation, the file might be deleted by the publisher.

E N D

Presentation Transcript

  1. REDOX Oxidation-Reduction Reactions

  2. Oxidation-Reduction Reactions • Electrons are transferred from 1 atom to another. • All single-replacement & combustion rxns are redox rxns.

  3. Oxidation = loss of electrons. LOSS of ELECTRONS = OXIDATION LEO

  4. Reduction = gain of electrons. GAIN of ELECTRONS = REDUCTION GER

  5. REDOX • Oxidation & Reduction are complementary. • They occur together & simultaneously or not at all.

  6. LEO GOES GER!!! Memorize! Oxidation of Cu

  7. Oxidation Numbers • In Ionic Compounds: the number of electrons lost or gained by an atom when it forms ions. Oxidation states of Vanadium

  8. Assigning Oxidation Numbers 8 RULES

  9. 8 Rules for Oxidation Numbers 1. of a free, uncombined element = 0. Na He O2 N2 S8 Cl2 P 2. of a monatomic ion = charge on ion. Ca+2 = +2. Cl-1 = -1. Al+3 = +3. Remember: Ions occur in ionic compounds: CaCl2, Al(NO3)3, etc.

  10. 8 Rules for Oxidation Numbers 3. CF4 Fluorine is always -1. 4. Hydrogen is nearly always +1, except when it’s bonded to a metal. Then it’s -1. LiH CaH2 NaH H2O, HNO3, H2SO4

  11. 8 Rules for Oxidation Numbers 5. Oxygen is nearly always -2 except when its • -Bonded to fluorine, where O is +2 OF2 • -In the peroxide ion, where O is -1. O22-

  12. 8 Rules for Oxidation Numbers 6. The sum of oxidation numbers in a neutral compound is 0. H2O CO2 NO SO3 7. The sum of oxidation numbers in a polyatomic ion = charge of the ion. Sum in SO42- = -2. Sum in NO3- = -1.

  13. 8 Rules for Oxidation Numbers 8. In covalent compounds, the oxidation number of the more electronegative atom is the negative charge it would have if it was an ion. *NH3: N = -3, H = +1. SiCl4: Si = +4, Cl = -1.

  14. KCl CaBr2 CO CO2 Al(NO3)3 Na3PO4 H2S NH4+1 SO3-2 Assign Oxidation Nos K = +1, Cl = -1 Ca = +2, Br = -1 C = +2, O = -2 C = +4, O = -2 Al = +3, O = -2, N = +5 Na = +1, O = -2, P = +5 H = +1, S = -2 N = -3, H = +1 S = +4, O = -2

  15. Electrons are Negative! • Why do we use the word “reduced” when electrons are gained? Look at how the oxidation number changes. For example, if Cl gains an electron it becomes Cl-1. The oxidation number decreased from 0 to -1. The oxidation number was reduced.

  16. Writing Equations • Even though oxidation & reduction occur together, we can write separate equations for each process. • Called Half-Reactions. • In order to balance a redox equation, we have to split the full equaton into half-reactions.

  17. Conservation of Mass • # of atoms of each type is the same on both sides of the equation. • Still holds for half-reactions. • Do this step first.

  18. Conservation of Charge • Total charge on LHS must equal total charge on RHS. • In the past, we usually had both sides neutral. (0 = 0.) • Note: Total charge can be nonzero. Just has to be equal on the 2 sides. • If not balanced, add electrons to whichever side is too positive.

  19. Reduction Half-Reactions Electrons are gained so they are like a reactant! • I2 + 2e- 2I- • O2 + 4e-  2O-2 • Half-reactions must demonstrate conservation of mass & conservation of charge. • # of atoms of each element on LHS equals “ “ “ “ “ “ “ RHS. • Total charge on LHS = Total charge on RHS

  20. Oxidation Half-reactions Electrons are lost so they appear on the product side! • K  K1+ + 1e- • Fe2+  Fe3+ + 1e- • Cu  Cu2+ + 2e- • Total Charge on LHS = Total Charge on RHS • # atoms LHS = # atoms RHS

  21. Identifying Half-Reactions • Reduction: electron term is on reactant side. • Oxidation: electron term is on product side.

  22. Vocabulary Interlude • Oxidizing Agent: Is itself reduced. Accepts electrons from something else – aids oxidation for another species. • Reducing Agent: Is itself oxidized. • Loses electrons to something else – aids reduction for another species.

  23. Figuring out what is what! • Given an unbalanced equation. • Goal: Balance it. • Procedure: • Assign oxidation numbers to everything • Split into half-reactions • Balance them separately • “Match” the electrons • Add them together

  24. What’s oxidized & what’s reduced? USE OIL RIG

  25. 4 3 2 1 0 -1 -2 -3 -4 2) And if you’re lucky you strike oil & it shoots up OIL RIG 1) You dig down with an oil rig

  26. Oxidizing & Reducing Agents • They are both ALWAYS on the reactant side. • Identify them by seeing how the oxidation numbers change. Mg + Cu2+ Mg2+ + Cu 4 3 2 1 0 -1 -2 -3 -4 0 +2 +2 0 OIL RIG Mg is oxidized, so Mg is the reducing agent!

  27. Oxidizing & Reducing Agents • What’s oxidized & what’s reduced: Ca + FeCl2 CaCl2 + Fe • Assign oxidation numbers • Figure out what increases & what decreases. +2, -1 0 +2, -1 0 4 3 2 1 0 -1 -2 -3 -4 OIL RIG Ca is oxidized; Fe2+ is reduced. Ca = reducing agent; FeCl2 = oxidizing agnt.

  28. Oxidizing & Reducing Agents 4, -1 2 FeBr3 + SnBr2 2 FeBr2 + SnBr4 Fe3+ is reduced to Fe2+ Sn2+ is oxidized to Sn4+ FeBr3 is the oxidizing agent. SnBr2 is the reducing agent. 3, -1 2, -1 2, -1

More Related